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Chapter 3: Acid–Base Titration Back To Chapter

3.3: Titration of a Weak Acid with a Strong Base

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Analytical Chemistry

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Titration of a Weak Acid with a Strong Base

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3.2: Titration of a Strong Acid with a Strong Base

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3.4: Titration of a Weak Base with a Strong Acid

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Titration calculations for weak acids and strong bases involve different approaches, depending on the primary reactant.

Initially, 50 mL of 0.1 M acetic acid has a pH of 2.87, calculated with the K_a and an ICE table.

Post titration with 0.1 M NaOH, the solution forms a buffer.

Adding 10 mL of NaOH creates 0.001 moles of acetate, leaving 0.004 moles of acetic acid. The resultant pH is 4.14, calculated using the Henderson-Hasselbalch equation.

Halfway through, the pH equals pK_a due to equal acetic acid and acetate ion concentration.

At the equivalence point, the addition of 50 mL of NaOH converts all acetic acid to acetate, resulting in a pH transition to basic. Using an ICE table and K_b for acetate ions, the pH is found to be 8.72.

Any further addition of NaOH will dictate the pH, as it's a stronger base than acetate. For instance, adding 70 mL of NaOH results in a final pH of 12.22.

3.3: Titration of a Weak Acid with a Strong Base

In titrating a weak acid with a strong base, different calculation methods are applied at various stages. Initially, the pH of a weak acid like acetic acid is calculated using its dissociation constant (K_a) and an ICE table. Upon addition of a strong base such as sodium hydroxide, a buffer forms, and its pH is determined using the Henderson-Hasselbalch equation. As more base is added and the titration reaches the halfway point, the pH becomes equal to the pK_a of the acid, indicating equal concentrations of the acid and its conjugate base. At the equivalence point, all the acid is converted to its conjugate base, and the pH is calculated using the base's dissociation constant (K_b) and an ICE table. Beyond the equivalence point, the pH is governed by the concentration of the excess strong base.

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